1. I. Types of Chemical Bonds
Ionic Bonding
Chemical bond made by attraction of oppositely charged atoms after the transfer of electrons between them
2 Na Cl Na+Cl- (ionic compound)
Metals lose electrons to become cations
Na loses 1 e- [Ne]3s [Ne]
Nonmetals gain electrons to become anions
Cl gains 1 e- [Ne]3s23p [Ar]
The +/- charge attraction in the NaCl molecule makes it
8.37 x J more stable than 2Na and Cl2 apart. (Coulomb’s law)
Ionic compounds stay bound together as solids, but dissolve in water to give separated cations and anions
NaCl(s) H2O(l) Na+(aq) Cl-(aq)
3. Ionic Lewis Structure:

2. B. Formation of Binary Ionic Compounds
1. Lattice Energy = DE when separated gaseous ions pack to form a solid
a. Exothermic process since attraction of ions is favorable
b. Lattice Energy has a negative sign
c. Calculate by summing steps: Example Li(s) + ½ F > LiF(s)
2. The Born-Haber Cycle:
C. Lattices and Coulomb’s Law
Larger charges (+2/-2) attract more
CaO has > 4x the lattice E of NaF
1. Li(s) ----> Li(g) DE = +161kJ/mol
2. Li(g) ----> Li+(g) + e- DE = +520kJ/mol
3. ½ F2(g) ----> F(g) DE = +77kJ/mol
4. F(g) + e > F-(g) DE = -328kJ/mol
5. Li+(g) + F-(g) ----> LiF(s) DE = -1047kJ/mol
Li(s) + ½ F2(g) ----> LiF(s) DE = -617kJ/mol

3. Making Mg2+ and O2- ions requires
More energy than making Na+ and F-
Attractive forces of MgO >> NaF
Lattice energy 4x greater, but overall
energy difference is small

4. Predicting Formulas NaCl CaO
D. Noble Gas Electron Configurations (NGEC)
All elements react by gaining, losing, or sharing electrons to achieve the electron configuration of a noble gas (filled electron shells)
Two nonmetals react: They share electrons to achieve NGEC.
a. H (1s1) + H (1s1) H2 (1s2 = [He])
b. 2 Cl ([Ne]3s23p5) Cl2 ([Ar])
A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.
a. Na loses 1 e- [Ne]3s [Ne]
Cl gainis 1 e- [Ne]3s23p [Ar]
Ca lose 2 e- [Ar]4s [Ar]
O gains 2 e- [He]2s22p [Ne]
NaCl
CaO

5. E. Use Periodic Table to predict charges and formulas
Duet Rule: H and He always react to have 2 e- in outer (n =1) shell
Octet Rule: All other main group elements react to have 8 e- in outer shell
Transition metals and heavier metals: Harder to predict
E. Use Periodic Table to predict charges and formulas
MgxOy AlxSy CaxCly
F. Ion sizes: More negative ions are larger, more positive are smaller
Larger attraction for fewer electrons contracts cations: Ca > Ca+ > Ca2+
Less attraction for more electrons enlarge anions: O2- > O- > O
O2> F > Na+ > Mg2+ > Al3+
Ions get larger down a group because larger shells are involved

6. Sizes of Ions Related to Positions of the Elements in the Periodic Table

7. G. Trends in Lattice Energies
Increased charge leads to larger lattice energies
Smaller ions lead to larger lattice energies
H. Properties of Ionic Compounds
High melting Points—strong attraction of full charges keeps ions together
Conduct Electricity when dissolved in water—ions dissolve into water

8. I. Covalent Bonding
Chemical bond formed by the energetically favorable sharing of electrons between 2 atoms
Like charges repel each other with a similar strength that opposite charges attract one another
+/+ proton/proton and -/- electron/electron interaction are unfavorable between any two bonded atoms
The +/- attraction of an ionic bond overcomes this repulsion
In covalent bonds, the atoms are neutral, so some other way of interaction must help overcome the +/+ and -/- repulsion
Sharing the electrons between them can create 2 favorable +/- interaction
The covalent bond will only form if the result is overall lower energy

9. In a covalent bond, the 2 atoms involved must come together until a minimum energy is found.
Far away, there is no favorable interaction
Closer together, the +/+ and -/- repulsions increase energy greatly
If the atoms can share their electrons at lower energy, they form a covalent bond with a specific Bond Length
Bond Energy = amount of energy required to pull them apart again
H2 (covalent) is favorable, H+H- (ionic) is not

10. J. Polar Covalent Bonds
Unequal sharing of the e- between two covalently bonded atoms
This type of bond is somewhere between Ionic and Covalent
A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

11. Electronegativity = attraction for shared electrons
Electronegativity increases from left to right on periodic table
Electronegativy decreases down a group
The most electronegative atom in a bond
has the electrons closest to it most of the time
Atoms with the same electronegativities
form purely covalent bonds
Atoms with different electronegativities
form polar covalent bonds

12. Bond Polarity and Dipole Moment
Molecules with +/- ends have a dipole moment
d tells us it is not a full charge
Arrow points to negative end
Complex molecule shapes:
Dipole moments can cancel out to give a nonpolar molecule

13. K. Partial Ionic Character of Covalent Bonds
Range of Bond Character
Few bonds are purely Ionic or Covalent
Most bonds have some character of both due to unequal electronegativity
Calculating the Percent Ionic Character
To compare molecules, we can calculate how Ionic they are
Ionic Character increases as electronegativity difference increases
No individual bonds are 100% Ionic (for the gas phase)
Everything > 50% Ionic is considered Ionic in the solid phase
Any melted compound that conducts electricity = ionic